Laws of Chemical Combination + Stoichiometry

1. Laws of Chemical Combination Chem1 SY 10-11 - kaaferrer

2. What makes compounds different?

3. Law of Constant Composition 1799 Joseph Proust a chemical compound contains the same elements in exactly the same proportions (ratios) by mass regardless of the size of the sample or source of the compound

4. For example, water always consists of oxygen and hydrogen atoms, and it is always 89 percent oxygen by mass and 11 percent hydrogen by mass

5. Does mass change during a chemical reaction?

6. Law of Conservation of Mass Lavoisier heated a measured amount of mercury to form the red oxide of mercury. He measured the amount of oxygen removed from the jar and the amount of red oxide formed. When the reaction was reversed, he found the original amounts of mercury and oxygen.

7. Law of Conservation of Mass 1744 Antoine Lavoisier matter can not be created or destroyed in ordinary chemical or physical changes. the mass of the reactants (starting materials) equals the mass of the products 2Mg (s) + O 2 (g) → 2MgO (s) 48.6 g 32.0 g 80.6 g

8. Example 10 grams of CaCO 3 on heating gave 4.4g of CO 2 and 5.6 of CaO . Show that these observations are in agreement with the law of conservation

9. Law of Multiple Proportions 1803 John Dalton States that when two elements combine to form more than one compound, the masses of one element which combine with a fixed mass of the other element are in ratios of small whole numbers

10. Example: Carbon monoxide (CO): 12 parts by mass of carbon combines with 16 parts by mass of oxygen. Carbon dioxide (CO 2 ): 12 parts by mass of carbon combines with 32 parts by mass of oxygen. Ratio of the masses of oxygen that combines with a fixed mass of carbon (12 parts) 16: 32 or 1: 2

11. Water has an oxygen-to-hydrogen mass ratio of 7.9:1. Hydrogen peroxide, another compound consisting of oxygen and hydrogen, has an oxygen-to-hydrogen mass ratio of 15.8:1. Ratio of the masses of oxygen that combines with a fixed mass of hydrogen is 7.9: 15.8 or 1: 2

12. Chemical Formula Relationships

13. How many number of ATOMS? Ca(NO) 3 NaNO 3 Ba (IO 3 ) 2

14. Molecular and Formula Weights The sum of all the atomic weights of each atom in its chemical formula. Formula mass = ∑ atomic masses in the formula unit

15. SEATWORK!

16. Percentage Composition from Formulas Percentage by mass contributed by each element in the substance. %element= (# of atoms of that element)(atomic mass of that element) ------------------------------------------------------------------------ formula mass of the compound

17. Stoichiometry The quantitative relationships between the substances involved in a chemical reaction, established by the equation for the reaction

18. Terminologies ATOM Smallest particle of an element MOLECULE Smallest unit particle of a pure substance ION An atom or group of bonded atoms with electrical charge because of an excess or deficiency of electrons ELEMENT Pure substance; CANNOT be broken down into 2 or more pure substances by chemical means COMPOUND Pure substance; CAN be broken down into 2 or more pure substances by chemical means

19. Atomic Mass Atomic Mass = used to numerically indicate the mass of an atom in its ground state, it is expressed in the non SI unit of u u = refers to unified atomic mass unit (formerly known as atomic mass unit or amu) 1 amu = 1/12 the mass of carbon-12 atom , therefore the mass of C-12 atom is made EQUAL to 12 amu   Carbon-12 atom is an isotope of carbon   1 amu = 1.66 x 10 -24 g

20. Atomic Mass Mass of e = 1/1800 of mass of p and n so it is negligible making the equation Atomic mass of 12 C = mass of p + mass of n subatomic particle charge Mass (amu) Neutron None 1.0087 ≈ 1 Proton Positive 1.0073 ≈ 1 Electron Negative 5.486 x 10 -4 ≈ 0

21. Atomic Mass vs. Average Atomic Mass Atomic Mass Average Atomic Mass For carbon it is 12 u not 12.01 u *Used to relate the fact that the numerical value assigned to each element in the periodic table reflects the average abundances of the atoms that compose a naturally occurring element *Related to isotopes *For carbon it is 12.01 u *Chemists often will use the term “atomic mass” when they are actually referring to average atomic mass of an atom.

22. Average Atomic Mass (calculation) Solve for the Average Atomic Mass of the element Boron Average atomic mass = ∑ (mass x percent abundance) Where ∑ means “sum” Isotope Mass (u) Percent Abundance 11 B 11.009305 80.1 10 B 10.012937 19.9

23. Relative Atomic Mass & Atomic Weight Have different meaning from atomic mass but synonymous with each other, although of different historical origins Relative Atomic Mass “Ratio” of the average mass of the atom to the unified atomic mass; dimensionless Atomic Weight The name Dalton used in the early 19 th century to numerically describe the weight of atoms relative to each other

24. The average mass of the molecules of a binary compound (non metal-non metal) Unit: u Ex.: The molecular mass of carbon dioxide gas, CO 2 is 28.01 u. Molecular Mass

25. The average mass of the molecules of an ionic compound (metal-non metal) Unit: u Ex.: The formula mass of barium chloride, BaCl 2 is 208.2 u. Formula Mass

26. The Mole Concept Used to describe the number of particles (atoms, molecules, etc.) that make up sample of matter “One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12.” 1 mol of any substance = 6.02 x 10 23 units of that substance, where 6.02 x 10 23 Avogadro’s Number, N

27. Molar Mass Mass of 1 mole of substance SI unit: g/mol Provides a bridge between mass and amount

28. Conversion GRAMS MOLES FORMULA UNITS Use molar mass Use N

29. Atomic mass Molar mass Unit: g/mol Unit: amu; u Atomic level Macroscopic level Numerically equivalent Comparison of Atomic and Molar Mass Total mass of p + and n o Mass of 1 mole

30. Empirical Formula and Molecular Formula

31. Shows the relative number of atom of each element in the compound EMPIRICAL FORMULA

32. A compound is found out to contain 20.0% carbon, 2.2% hydrogen, and 77.8% chlorine. Determine EF of the compound. Problem 1

33. Shows the actual number of atom of each element in the compound MOLECULAR FORMULA

34. MF = (EF)n where n is an integer therefore , n = M MF M EF where M is molar mass Relationship between EF & MF

35. A compound with the empirical formula C 2 H 5 has a molar mass of 58.12 g/mol. Find the molecular formula of the compound. Problem 2

36. The molar mass of a compound that is 54.6% carbon, 9.0% hydrogen and 36.4% oxygen is 88 g/mol. Find MF of the compound. Problem 3

37. The Mole One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12. 6.02 x 10 23  Avogadro’s Number

38. MOLAR Mass The mass in grams of one (1) mole of a substance. g/mol, grams per mole The molar mass of any substance in grams per mole is always numerically equal to the atomic, formula, or molecular mass of the substance in amu .

39. mole atoms gram Formula weight Avogadro’s number